science

1. THE HYDRONIUM ION
• The proton does not actually exist in aqueous
solution as a bare H+
ion.
• The proton exists as the hydronium ion (H3O+
).
• Consider the acid-base reaction:
HCO3
-
+ H2O  H3O+
+ CO3
2-
Here water acts as a base, producing the
hydronium ion as its conjugate acid. For
simplicity, we often just write this reaction as:
HCO3
-
 H+
+ CO3
2-

2. Conjugate Acid-Base pairs
• Generalized acid-base reaction:
HA + B  A + HB
• A is the conjugate base of HA, and HB is
the conjugate acid of B.
• More simply, HA  A-
+ H+
HA is the conjugate acid, A-
is the
conjugate base
• H2CO3  HCO3
-
+ H+

3. AMPHOTERIC SUBSTANCE
• Now consider the acid-base reaction:
NH3 + H2O  NH4
+
+ OH-
In this case, water acts as an acid, with OH-
its
conjugate base. Substances that can act as
either acids or bases are called amphoteric.
• Bicarbonate (HCO3
-
) is also an amphoteric
substance:
Acid: HCO3
-
+ H2O  H3O+
+ CO3
2-
Base: HCO3
-
+ H3O+
 H2O + H2CO3
0

4. Strong Acids/ Bases
• Strong Acids more readily release H+ into
water, they more fully dissociate
– H2SO4  2 H+
+ SO4
2-
• Strong Bases more readily release OH-
into water, they more fully dissociate
– NaOH  Na+
+ OH-
Strength DOES NOT EQUAL Concentration!

5. Acid-base Dissociation
• For any acid, describe it’s reaction in water:
– HxA + H2O  x H+
+ A-
+ H2O
– Describe this as an equilibrium expression, K (often
denotes KA or KB for acids or bases…)
• Strength of an acid or base is then related to the
dissociation constant  Big K, strong acid/base!
• pK = -log K  as before, lower pK=stronger
acid/base!
]
[
]
][
[
A
H
H
A
K
x
x



6. • LOTS of
reactions are
acid-base rxns
in the
environment!!
• HUGE effect on
solubility due to
this, most other
processes
Geochemical
Relevance?

7. Organic acids in natural waters
• Humic/nonhumic – designations for organic
fractions,
– Humics= refractory, acidic, dark, aromatic, large –
generally meaning an unspecified mix of organics
– Nonhumics – Carbohydrates, proteins, peptides,
amino acids, etc.
• Aquatic humics include humic and fulvic acids
(pKa>3.6) and humin which is more insoluble
• Soil fulvic acids also strongly complex metals
and can be an important control on metal
mobility

8. pH
• Commonly represented as a range between
0 and 14, and most natural waters are
between pH 4 and 9
• Remember that pH = - log [H+
]
– Can pH be negative?
– Of course!  pH -3  [H+
]=103
= 1000 molal?
– But what’s ?? Turns out to be quite small 
0.002 or so…
– How would you determine this??

9. pH
• pH electrodes are membrane ion-specific
electrodes
• Membrane is a silicate or chalcogenide
glass
• Monovalant cations in the glass lattice
interact with H+
in solution via an ion-
exchange reaction:
H+
+ Na+
Gl-
= Na+
+ H+
Gl-

10. The glass
• Corning 015 is 22% Na2O, 6% CaO, 72%
SiO2
• Glass must be hygroscopic – hydration of
the glass is critical for pH function
• The glass surface is predominantly H+
Gl-
(H+
on the glass) and the internal charge is
carried by Na+
glass
H+
Gl-
H+
Gl-
H+
Gl-
H+
Gl-
H+
Gl-
H+
Gl-
H+
Gl-
H+
Gl-
Na+
Gl-
Na+
Gl-
E1 E2
Analyte solution Reference solution

11. pH = - log {H+
}; glass membrane electrode
pH electrode has different
H+
activity than the solution
SCE // {H+
}= a1 / glass membrane/ {H+
}= a2, [Cl-
] = 0.1 M, AgCl (sat’d) / Ag
ref#1 // external analyte solution / Eb=E1-E2 / ref#2
E1 E2
H+
gradient across the glass; Na+
is
the charge carrier at the internal
dry part of the membrane
soln glass soln glass
H+
+ Na+
Gl-
 Na+
+ H+
Gl-

12. Values of NIST primary-standard
pH solutions from 0 to 60 o
C
pH = - log {H+
}
K = reference and
junction potentials

13. pKx?
• Why were there more than one pK for
those acids and bases??
• H3PO4  H+
+ H2PO4
-
pK1
• H2PO4
-
 H+
+ HPO4
2-
pK2
• HPO4
1-
 H+ + PO4
3-
pK3

14. BUFFERING
• When the pH is held ‘steady’ because of
the presence of a conjugate acid/base
pair, the system is said to be buffered
• In the environment, we must think about
more than just one conjugate acid/base
pairings in solution
• Many different acid/base pairs in solution,
minerals, gases, can act as buffers…

15. Henderson-Hasselbach Equation:
• When acid or base added to buffered system
with a pH near pK (remember that when pH=pK
HA and A- are equal), the pH will not change
much
• When the pH is further from the pK, additions of
acid or base will change the pH a lot
]
[
]
[
log
HA
A
pK
pH




16. Buffering example
• Let’s convince ourselves of what buffering
can do…
• Take a base-generating reaction:
– Albite + 2 H2O = 4 OH- + Na+
+ Al3+
+ 3 SiO2(aq)
– What happens to the pH of a solution containing
100 mM HCO3- which starts at pH 5??
– pK1 for H2CO3 = 6.35

17. • Think of albite dissolution as titrating OH-
into
solution – dissolve 0.05 mol albite = 0.2 mol OH-
• 0.2 mol OH-  pOH = 0.7, pH = 13.3 ??
• What about the buffer??
– Write the pH changes via the Henderson-Hasselbach
equation
• 0.1 mol H2CO3(aq), as the pH increases, some of this
starts turning into HCO3-
• After 12.5 mmoles albite react (50 mmoles OH-):
– pH=6.35+log (HCO3-
/H2CO3) = 6.35+log(50/50)
• After 20 mmoles albite react (80 mmoles OH-
):
– pH=6.35+log(80/20) = 6.35 + 0.6 = 6.95
]
[
]
[
log
HA
A
pK
pH



Greg Mon Oct 11 2004
0 10 20 30 40 50 60 70 80 90 100
5
5.5
6
6.5
7
7.5
8
8.5
Albite reacted (mmoles)
pH

18. Bjerrum Plots
• 2 D plots of species activity (y axis) and
pH (x axis)
• Useful to look at how conjugate acid-base
pairs for many different species behave as
pH changes
• At pH=pK the activity of the conjugate acid
and base are equal

19. pH
0 2 4 6 8 10 12 14
log
a
i
-12
-10
-8
-6
-4
-2
H2S0
HS-
S
2-
H+
OH
-
7.0 13.0
Bjerrum plot showing the activities of reduced sulfur species as a
function of pH for a value of total reduced sulfur of 10-3
mol L-1
.

20. pH
0 2 4 6 8 10 12 14
log
a
i
-8
-7
-6
-5
-4
-3
-2
6.35 10.33
H2CO3* HCO3
-
CO3
2-
H+
OH-
Common pH
range in nature
Bjerrum plot showing the activities of inorganic carbon species as a
function of pH for a value of total inorganic carbon of 10-3
mol L-1
.
In most natural waters, bicarbonate is the dominant carbonate species!

21. Titrations
• When we add acid or base to a solution
containing an ion which can by
protonated/deprotonated (i.e. it can accept
a H+
or OH-
), how does that affect the pH?
pH
0 2 4 6 8 10 12 14
lo
g
a
i
-8
-7
-6
-5
-4
-3
-2
6.35 10.33
H2CO3* HCO3
-
CO3
2-
H+
OH
-
Common pH
range in nature

22. Carbonate System Titration
• From low
pH to high
pH
Greg Wed Oct 06 2004
0 5 10 15 20 25 30 35 40 45 50
2
3
4
5
6
7
8
9
10
11
12
NaOHreacted (mmoles)
pH
Greg Wed Oct 06 2004
0 5 10 15 20 25 30 35 40 45 50
-16
-14
-12
-10
-8
-6
-4
-2
NaOHreacted (mmoles)
Some
species
w/
HCO
3
-
(log
activity)
CO2(aq) CO3
--
HCO3
-

23. Titrations  precipitate
Greg Wed Oct 06 2004
2 2.5 3 3.5 4 4.5 5 5.5 6 6.5 7
-6
-5.5
-5
-4.5
-4
-3.5
pH
Some
minerals
(log
moles)
Fe(OH)3(ppd)
Boehmite

24. BJERRUM PLOT - CARBONATE
• closed systems with a specified total carbonate
concentration. They plot the log of the concentrations of
various species in the system as a function of pH.
• The species in the CO2-H2O system: H2CO3*, HCO3
-
,
CO3
2-
, H+
, and OH-
.
• At each pK value, conjugate acid-base pairs have equal
concentrations.
• At pH < pK1, H2CO3* is predominant, and accounts for
nearly 100% of total carbonate.
• At pK1 < pH < pK2, HCO3
-
is predominant, and accounts for
nearly 100% of total carbonate.
• At pH > pK2, CO3
2-
is predominant.

25. pH
0 2 4 6 8 10 12 14
log
a
i
-8
-7
-6
-5
-4
-3
-2
6.35 10.33
H2CO3* HCO3
-
CO3
2-
H+
OH-
Common pH
range in nature
Bjerrum plot showing the activities of inorganic carbon species as a
function of pH for a value of total inorganic carbon of 10-3
mol L-1
.
In most natural waters, bicarbonate is the dominant carbonate species!